Chemistry 1009
Lab Manual
University of Louisiana
at Monroe
Department of Chemistry
2010
Version 2.0
Contents
Lab Session 1: Laboratory Safety Rules and Check In..................................................
Fire, Injury, Spills and Cleanliness ..............................................................................
Desk Assignment Sheet (Chemistry 1009) ..................................................................
Commonly Used Equipment (not in the desk drawer).................................................
1
2
4
5
Lab Session 2, Experiment 1: Introductory Exercises ...................................................
Report Form 1 ..............................................................................................................
10
17
Lab Session 3, Experiment 2: Oxygen ...........................................................................
Report Form 2 ..............................................................................................................
18
21
Lab Session 4, Experiment 3: Preparation of Sodium Chloride ....................................
Report Form 3 ..............................................................................................................
22
25
Lab Session 5, Experiment 4: Law of Definite Proportions ..........................................
Report Form 4 ...........................................................................................................
26
30
Lab Session 6, Experiment 5: Hydrogen and the Activity Series of Metals..................
Report Form 5 ..............................................................................................................
32
37
Lab Session 7, Experiment 6: Acid-Base Titration .......................................................
Report Form 6 ..............................................................................................................
38
40
Lab Session 8, Experiment 7: Antacids .........................................................................
Report Form 7 ..............................................................................................................
41
43
Lab Session 9, Experiment 8: Calorimetry, Heat of Reaction .......................................
Report Form 8 ..............................................................................................................
44
50
Lab Session 10, Experiment 9: Charles' Law ................................................................
Report Form 9 ..............................................................................................................
51
53
Lab Session 11, Experiment 10: Determination of the Molar Mass of Oxygen ............
Report Form 10 ............................................................................................................
54
57
Lab Session 12, Experiment 11: Valence, the Combining Capacity of Elements and
Molecules ........................................................................................................................
Report Form 11 ............................................................................................................
58
62
Lab Session 13, Makeup Experiment: Copper Analysis by Complexometric Titration
Report Form 12, Makeup Lab ......................................................................................
63
65
To the Students:
Each experiment write-up has blanks within the procedures in which you should enter
requested information (such as calculated results or other responses). Also provided at
the end of each experiment is a Report Form, on which the same entry blanks are
duplicated. At the end of a laboratory period, your instructor may simply inspect the
completed blanks in the body of the experiment to verify that you have completed the
work. Alternatively, s/he may have you copy the results neatly onto the Report Form to
be handed in at the end of the period.
ii
Lab Session 1: Laboratory Safety Rules and Check In
These rules are designed to ensure that all work done in the laboratory will be safe for you and
your fellow students:

Wear safety goggles at all times when you are in the
laboratory! Your instructor will specify what type of
goggles/glasses is required.
Figure 1.1

Bare skin must be minimized while in the laboratory by wearing clothing that covers
one’s feet, legs and body completely. Hence, closed shoes and full-length pants are
required, since broken glass and spilled chemicals are all too common on the floors of
chemistry laboratories. Sandals, flip-flops, skirts, shorts, three-quarter length pants, bare
midriffs, bare backs and bare shoulders are not allowed. Long-sleeve shirts are
recommended but not required. Long hair should be tied back. Hats are not allowed.

No horseplay, joking or playing is permitted in the lab. Failure to obey this rule is cause
for immediate expulsion from the lab.

No eating or drinking is permitted in the lab or in the prelab classroom.

No visitors are allowed in the lab. Your friends should visit with you before or after, but
not during the lab session.

Read all labels carefully. Never use materials
from unmarked bottles. Most bottles will have
short paragraphs on the hazards of a chemical.
Most will have a safety sticker showing whether
or not the chemical is poisonous or corrosive,
and indicating hazards and precautions. See
Figure 1.2 for a sample.

Report all accidents to your laboratory
instructor.

Report all spills or breakage to your laboratory
instructor. Clean spills promptly with materials
provided by your instructor.

Inquire about the location of eye-wash and
shower stations in your lab.

Locate spill control materials in your lab.

Locate all fire extinguishers in your lab.
1
Figure 1.2
•
Do not throw chemicals down the drain without first consulting your lab instructor.
Never put solid material in the sink or down the drain. Solids are to be disposed of in the
trash containers.
•
Read the safety chart in the lab. This chart contains useful information and recommends
safe laboratory procedures.
•
Do not pipet liquids by mouth. Use a bulb to siphon liquids into a pipet.
•
Use the fume hoods to pour noxious or irritating chemicals, and to run chemical reactions
that generate noxious or irritating products.
•
Never work in the laboratory unsupervised.
•
Read the corresponding chapters of this manual before coming to class. This will
familiarize you with any potential hazards that may exist or evolve during the exercise.
Pay particular attention to any information concerning handling or safety of particular
chemicals or solutions.
Fire
Your first responsibility is to get out of harm’s way and inform those around you and your lab
instructor of the situation. If possible, turn off all gas cocks and remove flammable materials
from the area. Usually a wet towel thrown over the fire will extinguish the blaze. Do not throw
water on the fire; use the fire extinguisher if possible. In case of fire on your clothing, WALK,
DO NOT RUN, to the nearest fire blanket and wrap it around you. Inform the instructor as soon
as possible about the incident.
Injury
Report all accidents to the laboratory instructor immediately. In all but the most trivial cases, a
visit to the infirmary is required. Chemical splashes may require the use of the shower located in
each room. Chemical injury requires immediate attention. Wash the area thoroughly with soap
and water and report to the lab instructor. If chemicals get in your eye, flush the eye with water
for several minutes and do not touch the eye. When there is a question of physical injury,
through chemical or mechanical means, modesty should not be a factor. Showers and eyewash
stations are there to be used.
Spills
All spills must be reported to the stockroom. All chemical spills are a potential hazard. For acid
spills, use sodium bicarbonate or spill material provided in a hood in each lab. For base spills,
use a weak acid and water to neutralize the base and clean up the spill. When cleaning a spill,
wear gloves, apron, and goggles. Always ventilate the area when cleaning a spill. Mercury
spills, including broken thermometers, are to be treated by stockroom personnel. Do not attempt
to clean mercury spills or handle mercury!
2
Cleanliness
It is important to keep the laboratory as clean as possible, for safety reasons as well as aesthetic
reasons. Each pair of students is responsible for their immediate desk area. Before leaving the
laboratory, students should make sure that the area of the laboratory bench near their assigned
drawer is clean and dry, that Bunsen burners and other shared equipment are put away in the
appropriate space, and that the trough to the sink is free of any solid material. Laboratory
instructors will check work areas before approving completion of the experiment.
Pairs of students will be assigned dates for which they are responsible for cleaning the reagent
shelves, balances and balance tables in the weighing room, and surfaces under the hoods. This
duty will be rotated so that each pair of students will be responsible for general laboratory
cleanliness at least once during the semester.
3
Desk Assignment Sheet (Chemistry 1009)
Please
Print
Name
ULM ID
Name
ULM ID
Section #
3
2
1
2
1
1
1
1
1
1
1
Room #
Desk #
Beaker, 100 or 150 ml
Beaker, 250 ml
Beaker, 400 ml
Bottles, one 4 oz
and one 8 oz, round
Wash Bottle
Clamp, Pinchcock
Clamp, Test-Tube
Crucible and Cover
Cylinder, Graduated 10 ml
Cylinder, Graduated, 100 ml
Evaporating Dish
2
1
1
1
2
1
1
1
10
1
1
1
Flask, Erlenmeyer, 125 ml
Flask, Erlenmeyer, 250 ml
Flask, Florence, 500 ml
Funnel, Narrow Stem
Watch Glass
Spatula
Combustion Spoon
Test Tube, 25 × 150
Test Tube, 15 × 125
Thermometer
Tongs
Triangle, Wire
See pictures of items on the following pages
Check In:
Signature
Signature
Checked in by
Date
I MUST WEAR SAFETY GOGGLES DURING EACH LAB SESSION.
Initials_____ Date________
Initials_____ Date________
Check Out:
Checked out by
Date
4
Beakers
Bottle
Clamp, Pinchcock
Wash Bottle
Clamp, Test tube
5
Crucible and Cover
Evaporating Dish
Cylinder, Graduated
Flask, Erlenmeyer
Flask, Florence
6
Glass, Watch
Spatula
Funnel, Narrow Stem
Tongs
Combustion
Spoon
Wire Triangle
Test Tubes
7
Commonly Used Equipment (not in the desk drawer)
Top Loading Balance
Test Tube Brush
Squeeze Bulb
Bunsen Burner
Dispensing Bottle
Ring Clamp
Burette Clamp
Ring Stand Clamp
Burette
Two-Prong Clamp
Screw Clamp
Three Prong Clamp
8
Eye Dropper
(Suction Pipette)
Fire Blanket
Fire Extinguisher
Plastic Ruler
Striker
One-hole Stopper
Two-hole Stopper
Tripod
Metal or Plastic Trough
Ring Stand
Glass
Pipette
Pneumatic Trough
9
Thermometer
Thistle
Tube
Wire Gauze
Lab Session 2, Experiment 1: Introductory Exercises
These experiments are meant to provide your first experience in the laboratory. If you have prior
laboratory experience, these exercises constitute good review. Conduct the experiments in
whatever order seems convenient. Keep all results to show to your instructor before you leave,
or if requested, to turn in on the separate Report Form.
Weighing
You must learn how to use laboratory balances. As always, the limit of readable precision of the
scale should be recorded.
When approaching the balance you will need the following:
1. The substance to be weighed;
2. Any container or holder for the sample while on the balance;
3. Any sample handling device such as a spatula; and
4. Your notebook and a pencil or pen to record your measurement.
Balances in this laboratory have a semi-automatic tare (an allowance for mass of the container or
holder). On an electronic balance, "tare equals zero" is set by depressing the bar, which is also
the on/off switch: up for off, down for on, down again to tare. These balances have an automatic
range selector that will change the readout precision automatically to ± 0.01 g when the gross
mass on the pan is over 35 g. The measurement precision for small masses is best if very light
containers are used, such as the glassine weighing paper for dry solid samples. The precision for
samples less than 30 g gross mass is ± 0.001 g.
10
1A Experiment: Mass
1. Select five pennies from your resources. Include a new shiny one if possible. Dates are
useful for keeping them in order.
2. Without taring, place a sheet of paper on the pan. Read and record its mass in the table on
this page. When possible, use the cover to protect the pan from air drafts to obtain higher
precision.
3. Add the pennies to the pan one at a time, reading the mass after each addition. Enter each in
the table below in the column entitled "Cumulative Mass." Keep the pennies in order.
Calculate the mass of each penny by subtracting; record the differences in the column
entitled "Mass by Difference."
4. Remove all the coins and weigh each individually, taring to zero. Record the mass of each
coin in the table in the column entitled "Direct Weighing."
5. Calculate the mean or average mass ( m ) of the pennies: m = (∑ m ) / n .
6. Calculate the absolute deviation ( d = m – m ) from the mean of each mass (direct weighing)
and then determine the Average Deviation: d = (∑ d ) / n .
Cumulative
Mass
g
Mass
by Difference
st
g
g
g
g
nd
g
g
g
g
rd
g
g
g
g
th
g
g
g
g
th
g
g
g
g
Average:
g
g
Source
Weighing Paper
1 penny
2 penny
3 penny
4 penny
5 penny
Direct
Weighing
Deviation
from Mean
Consider two hypotheses:
1. All Lincoln-head pennies are manufactured with equal mass (within ± 0.001 g), but their
various histories result in different masses when measured.
2. New Lincoln-head pennies are lighter than older ones.
Observation is complicated by the various histories of the pennies. Some typical problems in
chemistry are illustrated here. Most obviously, the state of corrosion of the pennies represents an
uncontrolled experimental variable which can be important. Had we used non-circulated coins,
we would have expected better precision. On the other hand, pennies may not be very uniform
even when new. Another question arises: How much experimental difference is sufficient and
how consistently must it be observed for us to consider two data sets, or groups of data sets as
distinctly different? This is an important question for which statistical methods provide answers.
Which hypothesis do you choose, and why?
11
Density
Density is defined as the ratio of the mass of a sample to its volume. Mass and volume are
extensive properties of matter -- properties that depend on the quantities of substances. Such
properties are not of themselves useful in characterizing substances.
Intensive properties, on the other hand, are useful in characterizing substances. Intensive
properties are often determined by ratioing two extensive properties measured at constant
temperature (T) and pressure (P). Density is an example of this kind of intensive property.
When measured under known conditions of T and P, density can be used to characterize
substances. Of course, two or more substances may have the same density, but for a given
substance there is only one density (at constant T and P). If you determine that a colorless liquid
has a density of 1.00 g/mL at 4º C and 1 atm, this does not prove the liquid is water. This fact is
simply one piece of evidence that the substance may be water.
1B Experiment: The Density of Water at Room Temperature
1. Collect a small beaker of deionized water and measure its temperature.
T = _______________ ºC
2. Place a clean dry 10 mL graduated cylinder on the balance and tare it to zero. Carefully
transfer nearly 10 mL (but less) of water into the cylinder, taking care not to splash water up
on the sides (read the mass to ± 0.001 g).
"Weight" = Mass of H2O = ________________ g
3. Carefully read the volume occupied by the water at the bottom of the meniscus holding the
cylinder at eye level. The volume should be read to ± 0.1 mL.
Volume of H2O = _________________ mL
4. Calculate the density of water at the current temperature, noting the number of significant
figures.
D = Mass of H2O/volume of H2O = ________________g/mL
5. Consult the Table of Water Densities (below) and calculate the percent error in your
determination according to the following formula.
%Error = [|DTAB–DEXP|÷DTAB]×100 = ________________ %Error
Table of Water Density at Various Temperatures
T, ºC
D, g/mL
T, ºC
D, g/mL
T, ºC
D, g/mL
15
0.9991
20
0.9982
25
0.9970
16
0.9989
21
0.9980
26
0.9967
17
0.9987
22
0.9978
27
0.9965
18
0.9986
23
0.9975
28
0.9962
19
0.9984
24
0.9973
29
0.9960
12
1C Experiment: The Density of a Metal
Different metals are furnished for this determination, either as cylindrical rods or as pellets or
shot. In this experiment, you will determine the mass of a metal sample and its volume, then
calculate its density.
1. Start with a 10 mL graduated cylinder about half filled with H2O. Read the volume of water
in the cylinder.
Volume of H2O = ________________ mL
2. Put a cylindrical metal rod or a sample of metal pellets or shot into the water so that the
entire sample is submerged. Read the volume occupied by the water and the metal. Obtain
the greatest change in the water level consistent with having all the metal sample submerged
with no bubbles adhered to the surface of the sample. You may have to thump or jostle the
cylinder to get rid of bubbles.
Volume of H2O + Volume of metal = ________________ mL
3. The difference in the two volumes is the volume of the metal sample.
Volume of H2O + Volume of metal = ________________ mL
– Volume of H2O = ________________ mL
= Volume of metal = ________________ mL
4. Thoroughly dry the metal pieces and weigh them. Use brown paper towels to absorb most of
the water, then a hair dryer if available.
"Weight" = Mass of metal = ________________ g
5. Calculate the density of the metal and pick the metal from the Table of Metal Densities
(below).
D = Mass of metal ÷ Volume of metal = ________________ g/mL
Metal name _____________
Metal symbol ______
Table of Metal Densities (g/mL)
Metal
D, g/mL
Metal
D, g/mL
Metal
D, g/mL
Aluminum
2.70
Iron
7.87
Tin
7.29
Antimony
6.62
Lead
11.34
Titanium
4.51
Bismuth
9.80
Magnesium
1.74
Tungsten
19.30
Copper
8.94
Molybdenum
10.22
Zinc
7.13
1D Experiment: Length
1. While you are thinking about the precision of reading a balance, select a wooden splint and
measure its length on the inch scale and the centimeter scale.
Length = ________________ inches.
Length = ________________ cm
2. Convert the length in inches to length in centimeters using the factor 2.54 cm/in. Pay
attention to the significant figures in your results.
Length in inches converted to length in centimeters = ________________ cm
13
Adjusting a Bunsen Burner
When selecting a burner, check to see that the gas needle valve on the bottom will close
completely. Also check to see that the barrel of the burner will screw in and out so that the air
supply to the flame can be controlled.
With the burner gas valve off and the
hose connected to the burner and the
bench gas cock (see Figure 2.1), turn
the bench gas cock fully on and check
for leaks around the burner with a
match. With the air vents closed, open
the burner gas valve and light the
flame. The flame should be yellow and
luminous.
Stick a test tube into the flame briefly.
You should observe a deposit of carbon
black on the tube. When hydrocarbons
such as methane [CH4 (natural gas)]
burns in too little air (oxygen), the
reaction is:
Figure 2.1
CH4 (gas) + O2 (gas) → C (sol) + 2 H2O (liq)
With a bit more air, the flame becomes hotter and blue, but carbon monoxide is formed:
2 CH4 (gas) + 3 O2 (gas) → 2 CO (gas) + 4 H2O (liq)
Now adjust the air supply -- you may also have to adjust the gas with the
burner valve -- until the flame resembles Figure 2.2. This is the hottest flame
and is characterized by a blue inverted cone shape within the flame that is the
so-called reducing flame. A little above the apex of the cone is the hottest
area in the flame, reaching temperatures around 1500ºC.
Towards the top of the flame, conditions are oxidizing (high temperatures,
excess O2). The well-adjusted flame completely converts methane and
oxygen to carbon dioxide and water:
CH4 (gas) + 2 O2 (gas) → CO2 (gas) + 2 H2O (liq)
14
Figure 2.2
1E Exercise: Understanding Flames
Take a wooden splint and hold it with its edge resting on the top of the burner. Notice how the
splint is burned only on the edges of the flame. The flame under the cone is relatively cool
(about 350 ºC). Higher in the flame, the splint ignites uniformly. Show the splint to your
instructor along with the rest of today's results.
To turn off the Bunsen burner, execute the lighting procedure in reverse. Shut the gas valve at
the base of the burner, then close the close the bench gas cock.
Reactions and Separations
A very common and useful type of reaction is the double displacement reaction (also called a
metathesis or exchange reaction). This occurs readily among ionic compounds in solution when
two ions that form an "insoluble" salt are mixed. When the clear (not necessarily colorless)
solutions are mixed, a solid forms that "falls out" of solution or precipitates. The solid that is
formed, called the precipitate, may be separated from the remaining solution, which still contains
the counter-ions of the two salts that formed the precipitate. Upon evaporation, the other more
soluble product is recovered.
1F Experiment: Reactions and Separations I
Equal volumes of Na2CO3 and CaCl2 solutions will be mixed. The concentrations of these
solutions are such that equal volumes contain exact reacting quantities. The reaction is:
CaCl2 (aq) + Na2CO3 (aq) → CaCO3 (sol) + 2 NaCl (aq)
After mixing 5mL of each solution, note the formation of a white solid. Set up a filter as shown
in Figure 2.3. Catch the filtrate (the solution which passes through the filter paper), which
should be clear. Keep the precipitate that has been caught on the filter paper to show to your
instructor. What is the formula of the precipitate caught by the filter paper?
_____________________________________________________________________________
Set up an evaporating dish on a steam bath as shown in Figure 2.4 and evaporate a portion of the
filtrate. Keep the residue in the evaporating dish to show to your instructor. Be sure you know
what these compounds are.
15
Figure 2.4
Figure 2.3
1G Exercise: Reactions and Separations II
Barium sulfate, BaSO4 is very insoluble. Suppose solutions of Ba(OH)2 and H2SO4 (aq) are
mixed. Complete the equation below:
______ Ba(OH)2
(aq)
+ ______ H2SO4 (aq) → _____________ + _____________
16
Report Form 1: Introductory Exercises Name___________________________
Source
Weighing Paper
Cumulative
Mass
g
Partner_____________Section #_____
Mass
Direct
Deviation
By Difference
Weighing
From Mean
st
g
g
g
g
nd
g
g
g
g
rd
g
g
g
g
th
g
g
g
g
th
g
g
g
g
Average:
g
g
1 penny
2 penny
3 penny
4 penny
5 penny
1B Experiment: The Density of Water at Room Temperature
1.
2.
3.
4.
5.
T=_______ ºC
"Weight" = Mass of H2O = ______ g
Volume of H2O = ______ mL
D = Mass of H2O/volume of H2O = ______g/mL
%Error = [|DTAB–DEXP|÷DTAB]×100 = ______ %Error
1C Experiment: The Density of a Metal
6. Volume of H2O = ______ mL
7. Volume of H2O + Volume of metal = ______ mL
8. Volume of metal = ______ mL
9. "Weight" = Mass of metal = ______ g
10. D = Mass of metal ÷ Volume of metal = ______ g/mL
Metal name _____________
Metal symbol ____
1D Experiment: Length
11. Length = ______ inches
Length = ______ cm
12. Length in inches converted to length in centimeters = ______ cm
1E Exercise: Understanding Flames
Show the splint to your instructor along with the rest of today's results.
1F Experiment: Reactions and Separations I
Keep the precipitate that has been caught on the filter paper to show to your instructor.
What is the formula of the precipitate caught by the filter paper? ________________________
Keep the residue in the evaporating dish to show to your instructor.
1G Exercise: Reactions and Separations II
______ Ba(OH)2
(aq)
+ ______ H2SO4 (aq) → _____________ + _____________
17
Lab Session 3, Experiment 2: Oxygen
Oxygen is the most abundant element on or near the surface of the earth. Oxygen is present in
large quantities in the air, water, rocks, and minerals. Oxygen combines with almost all other
elements to form compounds that are called oxides of these elements.
Oxygen is not highly active (or reactive) at room temperature, but at higher temperatures its
activity (or reactivity) increases markedly.
Oxygen may be prepared in a number of ways. Liquefied air is the source of commercial
quantities of oxygen. Electrolysis of water is sometimes used to obtain small quantities of
oxygen (and hydrogen). The most common laboratory method is the thermal decomposition of
oxygen-containing compounds. This does not imply that any compound that contains oxygen is
a ready source of oxygen.
Metal oxides that are soluble in or react with water give basic solutions. These oxides are called
"basic oxides" or "basic anhydrides". Oxides of nonmetals that are soluble in or react with water
give acidic solutions, and these oxides are called "acidic oxides" or "acid anhydrides".
In this lab session, you will explore several factors that influence the speed of a chemical
reaction. These are temperature, concentration, and use of a catalyst. A catalyst is a substance
that changes the speed of a reaction, but is not consumed and may be recovered at the end of the
reaction. Ideally then, a catalyst can be used over and over again.
Procedural Note: This is the first experiment in which you must acquire a sample of a solid
crystalline powder from a reagent bottle. Certain rules should be followed when securing
reagents from a reagent bottle to prevent contamination of the reagent. It is bad procedure (thus
it is forbidden) to insert a spatula into a reagent bottle. It is equally improper to return unused
portions of a chemical to a reagent bottle. Thus, when securing a solid reagent (such as BaO2,
PbO2, KClO3 or MnO2 in this experiment), you should take a beaker or watch glass with you and
pour the needed amount of solid reagent into the beaker or watch glass. Alternatively, you may
take a piece of filter paper or weighing paper to the reagent shelf and pour the solid on the paper.
Excess solid reagent should be disposed of rather than returned to the reagent bottle. If the solid
in the reagent bottle cannot easily be poured because it has acquired moisture and lumped or
hardened, ask your lab instructor for help.
The same rules apply to liquid reagents. When obtaining a liquid reagent (either a solution or a
pure liquid) from a reagent bottle, no objects (such as eyedroppers or pipettes) should be inserted
into the reagent bottle. Liquid reagents should be poured from the reagent bottle into an
appropriate container, or secured through the delivery/dispensing system provided. Excess
liquid should not be returned to the reagent bottle.
18
2A Experiment
1. Place BaO2 in a small test tube to a height of about 1 cm.
2. Using another small test tube, do the same with PbO2.
3. Heat each of the test tubes in the flame of your burner. Use the "glowing splint" test to show
that oxygen is produced.
2 BaO2 (sol) → 2 BaO(sol) + O2 (gas)
2 PbO2 (sol) → 2 PbO (sol) + O2 (gas)
4. Clean and dry your large test tube. Add
KClO3 to the test tube to a height of 2 - 3
cm; add, using a spatula, a pinch of
catalyst, MnO2. Mix well. Mount the test
tube to a ring stand as shown in Figure 3.1.
Check to make sure that the glass tube
through the rubber stopper fits tightly, that
the rubber stopper fits tightly in the test
tube, and that the rubber tube for delivery
to the pneumatic trough fits tightly on the
glass tube through the rubber stopper. The
rubber tubes may need to be secured by a
hose clamp or by wire windings. Tight fits
prevent oxygen loss to the atmosphere.
5. Prepare to collect oxygen by displacement
of water. You will need two 8 oz bottles of
oxygen, so before initiating the
decomposition of KClO3, fill, invert, and
position in the pneumatic trough both
bottles. Heat the test tube, gently at first, to
start the decomposition. After each bottle
is filled, cover the mouth of the bottle with
a watch glass, remove it from the trough,
and set it right side up on the desk.
Figure 3.1
2 KClO3 (sol) ⎯Catalyst
⎯⎯→ 2 KCl(sol) + 3 O2 (gas)
(a) What does the collection of oxygen by displacement of water indicate about the solubility
of O2 in water?
19
(b) List four physical properties of oxygen.
(c) What is the purpose of MnO2?
6. Add about 10 mL of water to each of the oxygen-filled bottles. Move the watch glass aside
slightly and lower a glowing splint into the bottle. Allow the combustion to continue for
several seconds. Remove the splint, cover the bottle, and swirl to mix the water and gas.
(d) What gas was produced by the combustion?
(e) Write the equation for the reaction between the gas produced and the water.
(f) Test the solution formed with litmus paper. What do you observe?
(g) Classify the oxide formed by the combustion as acidic or basic:
7. Under the hood, place a small amount of sulfur in a combustion spoon. Ignite the sulfur with
the burner flame. Lower the burning sulfur into the second oxygen bottle.
(h) Record your observations.
(i) What gas is produced by the combustion?
(j) Write the equation for the reaction between the gas and water.
(k) Test the solution in the bottle with litmus paper. What do you observe?
(l) Classify the oxide formed by the combustion of sulfur as acidic or basic:
20
Name__________________________________
Partner____________________Section #_____
Report Form 2: Oxygen
(a) Is O2 soluble or insoluble in water?
(b) List four physical properties of oxygen.
(c) What is the purpose of MnO2?
(d) What gas was produced by the combustion?
(e) Write the equation for the reaction between the gas produced and the water.
(f) Test the solution formed with litmus paper. What do you observe?
(g) Classify the oxide formed by the combustion as acidic or basic:
(h) Record your observations.
(i) What gas is produced by the combustion?
(j) Write the equation for the reaction between the gas and water.
(k) Test the solution in the bottle with litmus paper. What do you observe?
(l) Classify the oxide formed by the combustion of sulfur as acidic or basic:
21
Lab Session 4, Experiment 3: Preparation of Sodium
Chloride
Sodium chloride will be synthesized by reacting sodium bicarbonate with hydrochloric acid. The
reaction equation is shown below:
NaHCO3
+
HCl (aq)
→
NaCl (aq)
+
H2O (aq)
(sol)
84.01 g
+
CO2
(gas)
36.46 g
58.44 g
18.02 g
44.01 g
The quantitative interpretation of the reaction is as follows: 84.01 g (1 mol) of sodium
bicarbonate reacts with 36.46 g (1 mol) of hydrochloric acid to generate 58.44 g (1 mol) of salt,
18.02 g (1 mol) of water, and 44.01 g (or 1 mol) of carbon dioxide. Of course the starting
quantity of NaHCO3 may be more or less than 84.01 g, but a proportionate quantity of the
hydrochloric acid will be consumed, and proportionate quantities of the products will be formed.
For example, should 100.00 g of NaHCO3 react with excess acid, the theoretical mass of salt
produced would be calculated as follows:
58.44 g sodium chloride
84.01 g sodium bicarbonate
100.00 g sodium bicarbonate
=
69.55 g sodium
chloride
Note that when a bicarbonate is reacted with excess acid, the salt produced is the only substance
not readily volatile. That is, the unreacted acid, the water, and the carbon dioxide are easily
removed by heating.
22
3A Experiment
Chemicals needed: sodium bicarbonate, concentrated hydrochloric acid.
1. Accurately weigh an empty, dry evaporating dish, and record its mass in blanks (b) and (e) in
the table below.
2. Add to the dish about 5 g of NaHCO3 and weigh again. Record the mass in blank (a) in the
table below.
3. Add 5 to 6 mL of distilled water to the dish to wet the bicarbonate. Cover the dish with a
watch glass.
4. Move the watch glass aside slightly and add, in small portions, about 6 mL of concentrated
hydrochloric acid from a 10 mL graduated cylinder. These small portions of acid should be
added so that the acid runs down the inside wall of the evaporating dish. After the addition
of 6 mL of acid, continue adding acid only as long as CO2 (gas) continues to be evolved.
5. Remove the watch glass and evaporate to dryness over a water bath (the evaporating dish is
placed on top of a beaker containing boiling water).
6. Next, heat the dish on wire gauze with the burner for about 3 minutes. Allow the dish to cool
and weigh accurately. Again, heat the dish, cool, and weigh. Continue heating and weighing
until the dish reaches constant mass. Record this constant mass in blank (d) in the table
below.
(a)
(b)
(c)
(d)
(e)
(f)
g
g
g
g
g
g
Mass of dish and NaHCO3
Mass of empty dish
Mass of NaHCO3 [(a)–(b)]
Mass of dish and residue
Mass of empty dish (b)
Mass of NaCl residue [(d)–(e)]
7. In the reaction studied, what reactant was present in limiting quantity?
3B Calculations
1. Calculate the theoretical yield of NaCl.
58.44 g NaCl
84.01 g NaHCO3
_________ g NaHCO3
=
___________________g NaCl
2. Calculate the percentage yield of NaCl.
Actual yield × 100
Theoretical yield
=
(
(
)
)
23
× 100
=
___________________%
3C Exercises
In the following problems, show calculations:
3. What theoretical mass of NaCl would result from reacting 60.00 g of NaHCO3 with excess
HCl (aqueous)?
58.44 g NaCl
84.01 g NaHCO3
60.00 g NaHCO3
=
___________________g NaCl
4. What theoretical mass of NaCl would result from reacting 3.00 moles of NaHCO3 with
excess HCl (aqueous)?
3.00 mol NaHCO3
1 mol NaCl
1 mol NaHCO3
58.44 g NaCl
1 mol NaCl
= ________________g NaCl
5. What mass of CO2 would be generated along with 35.00 g of NaCl?
35.00 g NaCl
58.44 g NaCl
44.01 g CO2
=
___________________g CO2
24
Report Form 3: Preparation of
Sodium Chloride
Name_________________________________
Partner___________________Section #_____
3A Experiment
(a)
(b)
(c)
(d)
(e)
(f)
Mass of dish and NaHCO3
Mass of empty dish
Mass of NaHCO3 [(a)–(b)]
Mass of dish and residue
Mass of empty dish (b)
Mass of NaCl residue [(d)–(e)]
g
g
g
g
g
g
In the reaction studied, what reactant was present in limiting quantity?
3B Calculations
1. Theoretical yield of NaCl.
2. Percentage yield of NaCl.
3C Exercises
3. What theoretical mass of NaCl would result from reacting 60.00 g of NaHCO3 with excess
HCl (aqueous)?
4. What theoretical mass of NaCl would result from reacting 3.00 mol of NaHCO3 with excess
HCl (aqueous)?
5. What mass of CO2 would be generated along with 35.00 g of NaCl?
25
Lab Session 5, Experiment 4: Law of Definite Proportions
The law of definite proportions states that when two or more elements combine to form a given
compound, they do so in fixed proportions by mass. This is the same as saying the composition
of a compound is fixed. For example, sodium chloride contains 39.3% by mass sodium and
60.7% by mass chlorine.
In this experiment a sample of KClO3 will be decomposed thermally, and the oxygen produced
will be expressed as a percentage of the original mass of KClO3. Experimental results will then
be compared to the theoretical percentage of oxygen in KClO3.
4A Experiment: Decomposition of a Salt
1. Add approximately 0.1 g of MnO2 to a clean, dry crucible and weigh the crucible including
the catalyst. Enter the mass on line (b) below.
2. Add KClO3 to the crucible until it is about one-third full. Mix the KClO3 and MnO2
thoroughly. Weigh again and enter the mass on line (a) below.
3. Place the crucible on a wire triangle supported on a ring clamp. Adjust the ring clamp to a
height that allows the crucible to be heated with the flame of your burner.
4. Heat the crucible slowly at first. Then regulate heating to prevent boil-over from the
crucible. The objective is to completely decompose the chlorate to the chloride and oxygen;
therefore, heat for about ten minutes after apparent total decomposition. Heating should be
so vigorous as to cause the bottom of the crucible to glow red. Heat until all purple color is
gone and the contents are a uniform gray color.
5. Allow the crucible to cool. Weigh the crucible and enter the mass on line (d) below.
6. Compute the following: mass of oxygen, mass of KCl, experimental percentage of oxygen,
theoretical percentage of oxygen, and percent error in the experimental percentage of oxygen,
calculated as [|actual yield–theoretical yield|/theoretical yield]×100.
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
g
g
g
g
g
g
%
%
%
Mass of crucible with MnO2 and KClO3
Mass of crucible with MnO2
Mass of KClO3 [(a)–(b)]
Mass of crucible and contents after heating
Mass of oxygen [(a)–(d)]
Mass of KCl [(c)–(e)]
Experimental percentage of oxygen
Theoretical percentage of oxygen
Percent error in experimental percentage of oxygen
7. Now assume that the compound you decomposed has the formula KClOx. Note that "x" is
the ratio of moles of O to moles of KCl. Complete the following table.
(j) Experimental moles of O
(k) Experimental moles of KCl
(l) Calculated value of above ratio
(m) Formula of compound decomposed
26
mol
mol
Hydrates
Hydrates are substances formed when water combines chemically in definite proportions with a
salt. The ratio of water molecules to the ions of the salt is a constant. Hydrates are not mixtures.
The anhydrous (without water) form of the salt is produced when all the water of hydration is
lost. Some examples of hydrates are listed below:
Formula
(CaSO4)2 • H2O
CaSO4 • 2H2O
CuSO4 • 5H2O
MgSO4 • 7H2O
Na2CO3 • 10H2O
Names
calcium sulfate hemihydrate "plaster of paris"
calcium sulfate dihydrate "gypsum"
copper (II) sulfate pentahydrate "blue vitriol"
magnesium sulfate heptahydrate "epsom salt"
sodium carbonate decahydrate "washing soda"
The • in the formula indicates a kind of chemical bond that usually can be easily broken. For
example, copper (II) sulfate pentahydrate can be converted to anhydrous copper (II) sulfate by
heating:
CuSO4 • 5H2O (sol) → CuSO4 (sol) + 5H2O (gas).
2+
In CuSO4 • 5H2O the bonding involves four water molecules coordinatively bound to the Cu
ion in a square planar structure (Figure 5.1) and one molecule of water bound to the sulfate ion
by hydrogen bonds (Figure 5.2).
Figure 5.2
Figure 5.1
Loss of hydration water to the atmosphere is called efflorescence. The property of some salts to
collect moisture from the air and dissolve in it is called deliquescence. A compound is
hygroscopic if absorption of water from the atmosphere occurs without dissolution of the
compound.
27
4B Experiment: Composition of a Hydrate
In the following experiment, record all masses in the data table.
1. Weigh your evaporating dish.
2. Introduce about 5 g of pulverized hydrated CuSO4. Note the appearance and color of the
solid; weigh the dish and contents.
3. Place the evaporating dish on the wire gauze and heat slowly until the color disappears.
Prolonged heating may result in the decomposition of the anhydride. What color change
occurred?___________________________________________________________________
4. When cool, weigh the dish and anhydride.
5. From the data, calculate the following:
the percentage of water in the hydrate [(mass of H2O) ÷ (mass of hydrate)] ×100,
the number of moles of water [(mass of H2O) ÷ molar mass of H2O],
the number of moles of anhydride [mass of anhydride ÷ molar mass of anhydride], and
the formula of the hydrate CuSO4 • x H2O where x = [(moles of water) ÷ (moles of
anhydride)].
6. Add water a few drops at a time to convert the anhydride back to the hydrate. Do you notice
a change in temperature? (Check with a thermometer.)______________________________
7. What evidence of a chemical change did you observe?_______________________________
(a) Mass of crucible or dish and hydrate
(b) Mass of empty crucible or dish
(c) Mass of hydrate [(a) – (b)]
(d) Mass of crucible or dish and anhydride
(e) Mass of anhydride [(d) – (b)]
(f) Mass of water in hydrate [(c) – (e)]
(g) Percentage of water in hydrate [(f) ÷ (c)] × 100
(h) Moles of water [(f) ÷ molar mass of H2O]
(i) Moles of anhydride [(e) ÷ molar mass of anhydride]
(j) Moles of water per mole of anhydride [(h)÷(i)]
(k) Theoretical moles of water per mole of anhydride
(l) % error [(|(j) – (k)| ÷ (k)) × 100]
(m) Formula of hydrate
(n) Equation for decomposition of your hydrate
28
Formula of Copper Sulfate Hydrate
Experiment
Duplicate
g
g
g
g
g
g
g
g
g
g
g
g
%
%
mol
mol
mol
mol
4C Experiment: Properties of hydrates
1. Place crystals of calcium chloride, washing soda, and lithium chloride on a watch glass and
expose them to the air of the laboratory.
2. Note the appearance of each sample towards the end of the lab. Which crystals deliquesce?
Which crystals effloresce?
Explain why one sample gained water and why one lost water.
4D Experiment
1. Introduce one spatula of each of the following into separate test tubes: sodium chloride,
potassium dichromate, aluminum sulfate, cobalt chloride, and nickel sulfate. Record color
before heating in the table below.
2. Heat each sample in turn, holding the tube at an angle of about 30º so that any water will
condense in the cool part of the tube. Record color after heading in the table below.
3. When you heated salts in this experiment, which ones lost water?
4. Can the original color be restored by addition of a few drops of water to the cooled tube?
List any hydrates that changed color upon heating, but for which the color came back when
water was added.
Color Before Heating
(a)
(b)
(c)
(d)
(e)
Sodium Chloride
Potassium Dichromate
Aluminum Sulfate
Cobalt Chloride
Nickel Sulfate
29
Color After Heating
Report Form 4: Law of Definite
Proportions
Name_________________________________
Partner___________________Section #_____
4A Experiment: Decomposition of a Salt
g
g
g
g
g
g
%
%
%
mol
mol
(a) Mass of crucible with MnO2 and KClO3
(b) Mass of crucible with MnO2
(c) Mass of KClO3 [(a)–(b)]
(d) Mass crucible and contents after heating
(e) Mass of oxygen [(a)–(d)]
(f) Mass of KCl [(c)–(e)]
(g) Experimental percentage of oxygen
(h) Theoretical percentage of oxygen
(i) Percent error in experimental percentage of oxygen
(j) Experimental moles of O
(k) Experimental moles of KCl
(l) Calculated value of above ratio
(m) Formula of compound decomposed
4B Experiment: Composition of a Hydrate
What color change occurred?
Do you notice a change in temperature? (Check with a thermometer.)
What evidence of a chemical change did you observe?
Experiment
(a) Mass of crucible or dish and hydrate
(b) Mass of empty crucible or dish
(c) Mass of hydrate [(a) – (b)]
(d) Mass of crucible or dish and anhydride
(e) Mass of anhydride [(d) – (b)]
(f) Mass of water in hydrate [(c) – (e)]
(g) Percentage of water in hydrate [(f) ÷ (c)] × 100
(h) Moles of water [(f) ÷ molar mass of water]
(i) Moles of anhydride [(e) ÷ molar mass of anhydride]
(j) Moles of water per mole of anhydride [(h) ÷ (i)]
(k) Theoretical moles of water per mole of anhydride
(l) % error [(|(j) – (k)| ÷ (k)) × 100]
(m) Formula of hydrate
(n) Equation for decomposition of your hydrate
30
g
g
g
g
g
g
%
mol
mol
Duplicate
g
g
g
g
g
g
%
mol
mol
4C Experiment: Properties of hydrates
Which crystals deliquesce?
Which crystals effloresce?
Explain why one sample gained water and why one lost water.
4D Experiment
When you heated salts in this experiment, which ones lost water?
Can the original color be restored by addition of a few drops of water to the cooled tube?
List any hydrates that changed color upon heating, but for which the color came back when
water was added.
Color Before Heating
(a)
(b)
(c)
(d)
(e)
Sodium Chloride
Potassium Dichromate
Aluminum Sulfate
Cobalt Chloride
Nickel Sulfate
31
Color After Heating
Lab Session 6, Experiment 5: Hydrogen and the Activity
Series of Metals
An activity series is a listing of chemical species in order of increasing (or decreasing) reactivity.
Metals (be sure you can locate them on the periodic table) tend to react by losing one or more
electrons (per atom). Non-metals tend to gain electrons when they react. When a metal, or any
chemical species, loses electrons, it is said to be oxidized. The word "oxidized" is used by
chemists to refer to the process that usually happens when atoms react with oxygen. Metals are
considered more active since they are more easily oxidized. One way an activity series may be
determined is to observe the spontaneity of reaction of a set of metals with a single oxidizing
agent.
The oxidizing agent used in most of the experiments in this lab session will be hydrochloric acid
(aqueous hydrogen chloride). In fact, any strong acid would do because the chemical entity
+
actually taking the electrons from the metals is the hydronium ion, H3O (aq), which is present in
all strong aqueous acids. Frequently, this species is simply called the hydrogen ion, or even the
+
hydrated proton, H (aq), since in aqueous solution it is hydrated.
+
The product formed when H (aq) is the oxidizing agent is H2 (gas), neutral molecular hydrogen,
+
which can only be formed if H (aq) can take an electron from the metal.
The ion formed from the metal in the reaction depends on the chemical nature of the metal.
Periodic Groups IA, IIA, and IIIA are very consistent in this matter. Group IA metals form +1
ions, etc. Transition metals are not as predictable, but the charges on their cations are known
from experiment. This knowledge is imparted to you by the charge shown in parentheses after
these metals on the activity series furnished below.
32
33
Metal + O2 (1 atm) → oxide
Metal + acid → H2 (gas) + salt
Metal + hot water → H2 (gas) + oxide or hydroxide
Metal + cold H2O
→ H2 (gas) +
hydroxide
K
(K1+)
Ba (Ba2+)
Sr
(Sr2+)
Ca (Ca2+)
Na (Na1+)
Mg (Mg2+)
Al (Al3+)
Mn (Mn2+)
Zn (Zn2+)
Cr (Cr3+)
Fe (Fe2+,3+)
Cd (Cd2+)
Co (Co2+)
Ni (Ni2+)
Sn (Sn2+)
Pb (Pb2+)
H2
(H1+)
Sb (Sb3+)
As (As3+)
Bi
(Bi3+)
Cu (Cu2+)
Ag (Ag1+)
Pd (Pd2+)
Hg (Hg2+)
Pt
(Pt2+)
Au (Au3+)
Forms oxides
only indirectly
Oxides reduced
by heat alone
Oxides reduced by H2 (gas)
Oxides not reduced by H2 (gas)
Activity Series for Metals
5A Experiment
1. A single example of an acid reacting with a metal will be performed and the products noted.
2. Add 50 mL of concentrated HCl (12 M) to 50 mL of water to make 100 mL of 6 M HCl
(save excess for later).
3. Using a 125 mL Erlenmeyer flask and a two-hole
stopper fitted with a glass tube, assemble the
hydrogen generator as shown in Figure 6.1. Make
sure all fittings (thistle tube through the stopper,
glass tube through the stopper, and rubber tubing
to the glass tube) are tight. A hose clamp or wire
winding may be needed.
4. Place 5 g of granulated zinc in the flask, close it
and add 25 mL 6 M HCl through the thistle tube.
Be sure your thistle tube's lower end is submerged.
Copious bubbles should form on the metal surface
with release of heat.
5. Begin catching the effluent of this apparatus with a
small test tube by letting it bubble into the waterfilled, inverted tube under water in a pneumatic
trough.
6. Being very careful to take the tube away from the
Figure 6.1
generator in an inverted position, light the gas
inside with a small flame. A small explosion will
occur, making a barking sound. Continue this procedure periodically. Eventually, as the
oxygen of the air is swept out of the generator, the product gas, hydrogen, will burn evenly
with a nearly invisible flame. Make observations on the properties of hydrogen.
Density as compared to air ___________________________________
Solubility in water ___________________________________
Reactivity with oxygen (write an equation) ___________________________________
7. Now we turn our attention to the other product of the reaction. When you have finished
testing the hydrogen, set your generator in the hood, open it and allow it to run down. Allow
all the hydrogen to escape as it forms and to not build up. Later, each group of two students
should decant a few mL of the solution off the metal and evaporate it over a steam bath. The
solid residue is the salt zinc chloride. Complete the following reaction:
____ Zn (sol) + _____ HCl (aq) → _____________ + _____________
Be sure you can write a similar equation for each metal in our activity series that reacts with
acid.
34
5B Experiment
1. The relative activity of some metals with acid will be observed.
2. While your hydrogen generator from Experiment 6A is exhausting itself, pick up a small
sample of each of the following metals: granulated zinc, tin, lead, copper, iron, aluminum,
and magnesium. If granules are not available, magnesium ribbon and aluminum turnings or
wire will do. Place each metal in its own test tube. The smallest amount you can see is the
right amount. Using too much of an active metal may cause the reaction contents to bubble
out of the test tube.
3. To each in succession add 2 mL of 6 M HCl and observe the reaction.
4. When you have tested all the metals, list them in order of decreasing reactivity and compare
your results with the order given in the activity series furnished. Judge the activity by speed
of gas evolution and heat given off.
Most
Least
Active
Active
5C Experiment
A very active metal such as sodium reacts with cold water. Water is a considerably weaker
oxidizing agent than 6 M HCl. Your instructor may choose to demonstrate this experiment.
1. The sodium is placed in a beaker of water placed behind a safety shield. Notice that the
reaction is so exothermic that it will ignite the hydrogen released. If no ignition occurs, place
the sodium on a small piece of filter paper floated on the water. The residual molten sodium
oxide sphere has a tendency to burst into pieces some time after the reaction seems to be
over.
BE CAREFUL -- Be sure you have your eye protection on while watching this!
2. Test the solution produced with red litmus paper. In addition to H2 (gas), a strongly basic
metal hydroxide is produced.
3. Complete the following equation:
_____ Na (sol) + _____ H2O (liq) → _____________ + _____________
Be sure you can write a similar equation for each of the very active metals of the activity series
that will react with cold water. This experiment may be repeated with calcium metal, if
available.
35
5D Experiment
An active metal reacts with oxygen and also reacts with hot water. As shown on the activity
series, most metals will react with oxygen at one atmosphere (in air, the partial pressure of O2 is
only 0.2 atm) to produce an oxide. Less active metals, which do not react with cold water, will
react with hot water to release H2.
1. Hold a 3 inch piece of magnesium ribbon by one end with your tongs. Light the other end in
the burner flame. (Careful, do not look directly at the flame; it is very bright.) Now quickly
plunge the burning ribbon into a 400 mL beaker of boiling water. How can it continue to
burn without the oxygen of the air?______________________________________________
2. Complete the two following equations:
_____Mg (sol) + _____O2 (gas) → __________________________________
_____Mg (sol) + _____H2O (liq) → __________________________________
At an elevated temperature, Mg(OH)2 will dehydrate to MgO as shown in the activity series, but
here the sparingly soluble base, Mg(OH)2, may be detected in water by turning litmus blue.
Since either a metal hydroxide or oxide will react with an acid to form a salt (and water, in the
case of hydroxide), the salt will always be a product when metal oxidation is done by acid, even
if the most active metals are involved.
5E Experiment
Apparently, activity is affected by conditions of concentration and temperature.
Part 1: More concentrated acid is a stronger oxidizing agent than dilute acid.
1. Put 1 mL of 6 M HCl in a 10 mL graduated cylinder and add 9 mL of water, then mix.
2. Place two similarly sized pinches of granulated zinc in two test tubes. To one add 6 M HCl
and to the other the dilute HCl.
3. What do you observe?
Part 2: If the metal is more finely subdivided, more contact between metal and acid occurs.
1. Place a pinch of zinc dust in a test tube and add 6 M HCl.
2. Compare this reaction to that of granulated zinc treated with 6 M HCl.
Part 3: The reaction rate activity should be enhanced at increased temperature.
1. Take a pinch of granulated tin and add 2 mL of 6 M HCl.
2. After noting the activity, heat carefully in a hot water bath.
3. What do you observe?
36
Report Form 5: Hydrogen and the Name_________________________________
Partner___________________Section #_____
Activity Series of Metals
5A Experiment
Density of H2 gas as compared to air ___________________________________
Solubility of H2 gas in water ___________________________________
____ H2 (gas) + _____ O2 (gas) → _____________
____ Zn (sol) + _____ HCl (aq) → _____________ + _____________
5B Experiment
Most
Active
Least
Active
5C Experiment
_____ Na (sol) + _____ H2O (liq) → _____________ + _____________
5D Experiment
_____Mg (sol) + _____O2 (gas) → __________________________________
_____Mg (sol) + _____H2O (liq) → __________________________________
5E Experiment
Part 1: Observation
Part 2: Observation
Part 3: Observation
37
Lab Session 7, Experiment 6: Acid-Base Titration
Molarity (M) and normality (N) are two means of expressing solute concentrations. Molarity is
defined as moles of solute per liter of solution, and in a similar way, normality is defined as
equivalents (eq) of solute per liter of solution. From these definitions we may write the
following:
M=
moles , or by rearrangement,
L
N=
N=
eq
L
, or by rearrangement,
millequivalents (or meq)
mL
M × L = moles
N × L = eq
, or by rearrangement,
N × mL = meq
In this experiment, we will consider only acid-base reactions. Your instructor will define the
quantity of an acid that is an equivalent and the quantity of a base that is an equivalent. Let us
stress here that one eq (or meq) of an acid reacts exactly with one eq (or meq) or a base.
Whatever the number of eq (or meq) of acid, the same number of eq (or meq) of base will be
consumed in the reaction. Thus,
meq acid = meq base
mL acid × N acid = mL base × N base
Assume that 20.00 mL of 0.3000 N acid requires 30.00 mL of base. Then,
Nbase =
mLacid × Nacid
mLbase
=
20.00 × 0.3000
30.00
= 0.2000 N
Our assumed case may be stated as follows: The titration of 20.00 mL of 0.3000 N acid requires
30.00 mL of 0.2000 N base.
Titration is the precise measurement of the volume of one reagent required to react with a mass
or volume of another reagent. As in the titration described above, the solution of base would be
added from a burette to the acid until the acid is just neutralized. You will use the indicator
phenolphthalein, which is colorless in acid and pink in base, to indicate when the acid is just
neutralized. Two or three drops of phenolphthalein are added to the acid before the titration is
begun. This colorless solution turns pink when, as a result of base addition from the burette, the
solution changes from acidic to barely basic. When the color change appears, the burette is read
to obtain the volume of base added.
38
6A Experiment
In this experiment, you will use a standardized NaOH solution to titrate against an aqueous acetic
acid unknown in order to determine the normality of the acetic acid solution.
1. You will be provided with three chemicals:
(1) An NaOH standard solution of known normality (approximately 0.21 N). The exact
normality should be recorded in the data table.
(2) Acetic acid, CH3CO2H, of unknown normality. This is to be prepared by the stockroom
personnel by diluting 12 mL of concentrated acetic acid to one liter.
(3) Phenolphthalein in a dropper bottle.
2. Clean and empty the burette, rinse with and then fill with the acetic acid.
3. Draw off three 15.00 mL samples of acetic acid into three Erlenmeyer flasks. You may use
either size flask, 125 mL and/or 250 mL.
4. Add 2-3 drops of indicator to each sample.
5. Clean and empty the burette, rinse with and then fill with the NaOH solution.
6. Titrate each acid sample to the phenolphthalein end-point. Complete the data table below for
each of the three titrations.
1
Titrations
2
3
Nbase
mLacid
mLbase
Nacid
% deviation
Averages
Nacid
% deviation
6B Exercise
1. Calculate the grams of acetic acid in one liter of solution.
2. The density of the acetic acid solution is 1.00 g/mL. Calculate the percentage of acetic acid
in the solution.
3. Give the equations for the neutralization of HCl (aq) and CH3CO2H (aq) with NaOH (aq).
Be sure to clean your burette and rinse it with deionized water before you return it to the
hood.
39
Report Form 6: Acid-Base Titration Name________________________________
Partner__________________Section #_____
6A Experiment
Titrations
2
1
3
Nbase
mLacid
mLbase
Nacid
% deviation
Averages
Nacid
% deviation
6B Exercise
1. Grams of acetic acid in one liter of solution
g
2. Percentage of acetic acid in the solution
%
3. Balanced Neutralization equations
_____ HCl(aq) + _____ NaOH(aq) → _____ H2O(aq) + _______________(aq)
_____ CH3CO2H(aq) + _____ NaOH(aq) → _____ H2O(aq) + ______________(aq)
Note: In the above table, if N1, N2 and N3 are the values of Nacid for the individual titrations,
Average Nacid =
N1 + N2 + N3
3
Defining Navg = average Nacid, the % deviations for the individual titrations are
% D1 =
|N1 – Navg|
Navg
X 100,
% D2 =
|N2 – Navg|
Navg
The average % deviation, therefore, is
Average % deviation =
%D1 + %D2 + %D3
3
40
X 100,
% D3 =
|N3 – Navg|
Navg
X 100
Lab Session 8, Experiment 7: Antacids
The stomach secretes hydrochloric acid, which is necessary for digestion of food. However, for
a number of reasons the stomach may secrete excess hydrochloric acid that causes discomfort. A
number of commercial antacids are available to neutralize the excess acid. It should be
remembered that antacids should be consumed in limited amounts in order that the stomach not
be made alkaline. In this experiment, several brands of antacid tablets will be compared for their
capacity to neutralize hydrochloric acid.
7A Experiment
In this experiment, you will add an antacid tablet to a known quantity of 0.50 M HCl that
exceeds the amount required for reaction with the tablet. Standardized NaOH solution will be
added to neutralize the excess acid. The millimoles of acid neutralized by the antacid will be the
initial number of millimoles of acid minus the number of millimoles of excess acid, which is the
number of millimoles of NaOH required.
Chemicals Required: Antacid tablets
0.50 M HCl
NaOH solution
Phenolphthalein indicator
CAUTION:
DO NOT TASTE THE ANTACID
TABLETS USED IN THIS LAB!
1. Dissolve the antacid tablet in 50.00 mL of the 0.50 M HCl solution. Heat very gently and
swirl the container if necessary to dissolve the tablet. A small part of the tablet may not
dissolve.
2. After dissolving as much of the tablet as possible, add 3 drops of phenolphthalein solution.
Since an excess of aqueous HCl solution has been added, the indicator should be colorless. If
it is not colorless but rather turns red, add 10.0 mL more of the HCl solution, and note that
the total volume of acid is 50.00 mL + 10.00 mL = 60.00 mL.
3. Titrate to the end point with the standardized NaOH solution. The molarity of the NaOH
solution will be given on the label of the reagent bottle.
(a) Mass of antacid tested
g
(b) Milliliters of HCl used
mL
(c) Molarity of HCl
mol/L
mmol
(d) Millimoles of HCl used = (mLHCl × MHCl)
(e) Milliliters of NaOH used to neutralize excess acid
mL
(f) Molarity of NaOH
mol/L
mmol
(g) Millimoles of NaOH used = (mLNaOH × MNaOH)
(h) Millimoles of HCl (total used)
mmol
(i) Millimoles of excess HCl*
mmol
(j) Millimoles of HCl neutralized by antacid [(h)–(i)]
mmol
(k) Mass of CaCO3
g
(l) % CaCO3 in tablet
%
*Millimoles of excess acid = millimoles of base required to neutralize excess acid.
41
Different types of antacid tables may be studied in your class. If so, you are encouraged to
compare the results.
7B Exercise
1. The active ingredient in one brand of antacid is CaCO3.
(a) Write the equation for the reaction between HCl (aq) and CaCO3 (sol).
(b) How many millimoles of HCl (aq) are neutralized by 0.35 g of CaCO3 (sol)?
mmol
2. Which household substance is frequently used as an antacid?
(a) Write the equation for the reaction between this substance and HCl (aq).
(b) How many millimoles of HCl (aq) are neutralized by 5.0 g of this substance?
3. Suggest an explanation for why an antacid tablet may not be completely soluble in acid
solution:
42
mmol
Report Form 7: Antacids
Name__________________________________
Partner____________________Section #_____
7A Experiment
(a) Mass of antacid tested
g
(b) Milliliters of HCl used
mL
(c) Molarity of HCl
mol/L
mmol
(d) Millimoles of HCl used = (mLHCl × MHCl)
(e) Milliliters of NaOH used to neutralize excess acid
mL
(f) Molarity of NaOH
mol/L
mmol
(g) Millimoles of NaOH used = (mLNaOH × MNaOH)
(h) Millimoles of HCl (total used)
mmol
(i) Millimoles of excess HCl*
mmol
(j) Millimoles of HCl neutralized by antacid [(h)–(i)]
mmol
(k) Mass of CaCO3
g
(l) % CaCO3 in tablet
%
*Millimoles of excess acid = millimoles of base required to neutralize excess acid
7B Exercise
1. The active ingredient in one brand of antacid is CaCO3.
(a) Write the equation for the reaction between HCl (aq) and CaCO3 (sol).
(b) How many millimoles of HCl (aq) are neutralized by 0.35 g of CaCO3 (sol)?
mmol
2. Which household substance is frequently used as an antacid?
(a) Write the equation for the reaction between this substance and HCl (aq).
(b) How many millimoles of HCl (aq) are neutralized by 5.0 g of this substance?
3. Suggest an explanation for why an antacid tablet may not be completely soluble in acid
solution:
43
mmol
Lab Session 9, Experiment 8: Calorimetry, Heat of Reaction
Specific heat is an intensive property of a single phase (solid, liquid or gas) sample that describes
how the temperature of the sample changes as it either absorbs or loses heat energy. Specific heat is
generally a function of temperature, but, to a good approximation, it can be treated as being constant
for a single phase over a moderate temperature range. The table below lists the specific heats at 25oC
of liquid water and selected metal solids.
The Zeroth Law of Thermodynamics states: “If two samples of matter, initially at different
temperatures (TH and TC, respectively), are placed in thermal contact, heat will be lost by the hotter
sample (TH) and gained by the cooler one (TC). This exchange of heat will take place until both
samples achieve the same final temperature, TF, so that
Substance
Specific Heat, J/(goC) TH > TF > TC.” The First Law of Thermodynamics
states: “During heat exchange (or, as we shall see,
water (liquid)
4.184
during a chemical reaction), heat is neither created nor
aluminum (solid)
0.901
destroyed.” Thus, heat lost equals heat gained, if the
chromium (solid)
0.448
heat exchange can be sufficiently insulated from the
copper (solid)
0.386
surroundings so that not very much heat escapes to the
iron (solid)
0.450
surroundings. A vessel that provides adequate heat
lead (solid)
0.129
insulation from its surroundings, and in which
nickel (solid)
0.443
temperature changes are measured in order to
tin (solid)
0.217
determine specific heats (or heats of chemical
zinc (solid)
0.386
reactions) is called a calorimeter. A perfect
calorimeter absorbs no heat from the solution that it
contains, nor loses any heat to the surroundings. No calorimeter is perfect, however. A typical
calorimeter used in freshman chemistry labs is made of two nested Styrofoam cups, and looks like the
one shown in Figure 9.1.
If we place mmetal grams of a hot
metal, at a Celsius temperature TH,
into a nearly perfect calorimeter,
along with mwater grams of water, at a
lower Celsius temperature TC, then
the metal and the water will thermally
equilibrate at a final intermediate
Celsius temperature, TF.
Since the heat lost by the metal is
equal to the heat gained by the water
within the calorimeter, the following
equation holds (m is in the units g, T
is in the units °C and sp ht is in the
units J/g °C):
Figure 9.1
[(mmetal) × (TH – TF) × (sp htmetal)] = [(mwater) × (TF – TC) × (sp htwater)] .
44
We can possibly identify an unknown metal by determining its specific heat using the preceding heat
exchange formula with the specific heat of water taken to be 4.184 J/g oC. From the table above, it is
obvious that we cannot distinguish between copper and zinc, or between chromium and nickel, on the
basis of specific heat alone. Additional information is needed, such as color. (Copper has the
characteristic "copper" color, whereas zinc is gray.) Rearranging the heat exchange formula and
inserting the specific heat of water gives
sp htmetal =
(m water )(TF − TC )(4.184 J/g°C)
.
(m metal )(TH − TF )
A major source of error in the preceding example is the assumption that no heat went into warming
the calorimeter vessel and that no heat was lost to the surroundings. We can take this into account by
defining the heat capacity of the calorimeter C, which has the units J/oC. By making a similar heat
exchange measurement in which we mix mH grams of hot water, initially at temperature TH, with mC
grams of water initially at a lower temperature TC (perhaps at or near room temperature). The value
of the calorimeter’s heat capacity reflects the calorimeter’s efficiency. (A more efficient calorimeter
has a smaller value of C.) With this correction for the calorimeter and surroundings, the heat
exchange relationship becomes:
heat lost by the hot water = heat gained by the cold water
+ heat gained by the calorimeter and surroundings,
or
[(mH)(TH –TF)(4.184 J/g °C)] = [(mC)(TF – TC)(4.184 J/g °C) + (C)(TF – TC)] .
Rearranging, we have
C =
=
(m H )(TH - TF )(4.184 J/g°C) - (m C )(TF - TC )(4.184 J/g°C)
(TF − TC )
(m )(T - T ) - (m C )(TF - TC )
](4.184 J/g °C).
[ H H F
(TF − TC )
This heat capacity is most reliable for the temperature range in which it was determined (i.e, for
temperature changes that occur just above room temperature in this case). The heat which is absorbed
and thereby increases the temperature of the water and the calorimeter can be generated by an
exothermic chemical reaction occuring within the calorimeter. For example, if we mix equimolar
amounts of an aqueous strong acid (e.g., HCl (aq)) with an aqueous strong base (e.g., NaOH (aq)), the
neutralization reaction produces heat, which causes the temperature in the calorimeter to increase.
If we mix 50.0 mL of a 2.00 M aqueous HCl solution with 50.0 mL of a 2.00 M aqueous NaOH
solution, the neutralization reaction HCl (aq) + NaOH (aq) Æ H2O (aq) + NaCl (aq) will produce a
measurable change in temperature. After complete reaction, the calorimeter will contain a dilute
solution of sodium chloride. We can assume that the specific heat is still that of water, namely
45
4.184 J/g oC, without introducing much error. We can also assume that the density of the dilute salt
solution is that of water, 1.00 g/mL. The heat exchange equation in this case is:
heat given off by the reaction = heat gained by the water
+ heat gained by the calorimeter and surroundings.
If the calorimeter, the HCl solution, and the NaOH solution all start at the same initial temperature TI
and warm to a final temperature TF, then
heat given off by the reaction = (mwater)(TF – TI)(4.184 J/g °C) + (C) (TF – TI)
= [(Vwater in mL)(1.00 g/mL)(4.184 J/g °C) + (C)](TF – TI).
(Note: Vwater = 100 mL)
At constant pressure, the heat given off by the reaction is equal to the change in enthalpy (∆H) for the
mol ⎛ 10 - 3 L ⎞
⎟ (50.0 mL) = 0.100 mol of HCl = 0.100 mol
reaction. Since the reaction involves (2.00
)⎜
L ⎜⎝ 1 mL ⎟⎠
of NaOH, we can calculate the change in enthalpy for the acid/base reaction as follows (the minus
sign is inserted to indicate that the reaction is exothermic):
⎛ - (heat) J ⎞ ⎛ 1 kJ ⎞ ⎛ - (heat) ⎞
∆H = ⎜
⎟⎜ 3 ⎟= ⎜
⎟ in kJ/mol.
⎝ 0.100 mol ⎠ ⎝ 10 J ⎠ ⎝ 100 ⎠
46
8A Experiment: Determination of Calorimeter Constant
1. Obtain or assemble a calorimeter as shown in Figure 9. The experiment will require two
thermometers, one for the calorimeter and one for the heated water.
2. Using a graduated cylinder, measure 50.0 mL of water and pour it into the calorimeter.
Measure an additional 50.0 mL of water and pout it into a clean, previously dried beaker.
3. One lab partner should stir the calorimeter contents for at least 5 minutes and then record the
temperature inside the calorimeter as TC.
4. Meanwhile, the other lab partner should heat (bunsen burner) and stir the water in the beaker
until it reaches a temperature of 55-60 oC. After removing the burner, stir and record the exact
temperature of the water in the beaker as TH.
5. At that point, pour the hot water into the calorimeter, replace the top and stir the contents well
while recording the temperature at 15 second intervals. The highest temperature should occur
within 2 minutes. Record this maximum temperature as TF.
6. Calculate the heat capacity C of the calorimeter in J/oC using the formula given in the
preceding section with mH = mC = (50.0 mL)(1.00 g/mL) = 50.0 g.
7. Repeat steps 1 through 6 in order to make a second determination of C. Reverse the
thermometers for the second determination of C.
First Determination of C
TC =
°C
TH =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
C=
J/°C
47
Second Determination of C
TC =
°C
TH =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
C=
J/°C
Average C =
J/°C
8B Experiment: Determination of Reaction Enthalpy
1. Dry the calorimeter and reassemble it.
2. Using a clean, dry graduated cylinder, measure 50.0 mL of a 2.0 M aqueous NaOH solution
and pour it into the calorimeter. (Caution: NaOH is corrosive and, if spilled on the skin,
should be washed away immediately with copious amounts of water.)
3. Using a clean, dry graduated cylinder, measure 50.0 mL of a 2.0 M aqueous HCl solution and
pour it into a clean, dry beaker. (Caution: HCl is corrosive and, if spilled on the skin, should
be washed away immediately with copious amounts of water.)
4. Stir the two solutions (the NaOH (aq) in the calorimeter, and the HCl(aq) in the beaker) until
they exhibit the same temperature. Record this temperature as TI.
5. Pour the HCl solution into the calorimeter, replace the top, and stir while recording the
temperature at 15 second intervals.
6. Record the highest temperature reached as TF.
7. Use these temperatures and the value of C determined for the calorimeter in the equations
presented in the previous section to calculate the heat evolved and ∆H of the reaction.
(# of moles = (0.0500 L)(2.0 mol/L) = 0.10 mol. ∆H = heat/moles.)
8. Repeat steps 1 through 7 with the thermometers reversed in order to make a second
determination of the heat evolved and of ∆H. (∆H should equal – 55.8 kJ/mol.)
First Determination of ∆H
TI =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
heat =
J
∆H =
kJ/mol
Second Determination of ∆H
TI =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
heat =
J
∆H =
kJ/mol
Average ∆H =
kJ/mol
% error = [(|∆H| - |55.8|)/55.8] x 100 =
%
48
8C Exercises
1. Suppose that one were to mix 30.0 g of aluminum pellets, originally at 97.0 oC, with 100
grams of water, originally at 23.0 oC, in a perfect calorimeter. What will be the
equilibrium temperature TF in the calorimeter? TF = ____________ oC.
2. Repeat the calculation of Exercise 1, but in the calorimeter you used in lab rather than in a
perfect calorimeter.
TF = ____________ oC.
3. When 100 g of an unknown metal at 98.0 oC were mixed in a perfect calorimeter with
50.0 g of water at 22.0 oC, the final temperature TF was observed to be 26.4 oC.
Presuming that the metal is one of those listed in the table of specific heats given above,
which one is it?
Unknown metal = _______________.
49
Report Form 8: Calorimetry,
Heat of Reaction
Name___________________________________
Partner_____________________Section #_____
8A Experiment: Determination of Calorimeter Constant
First Determination of C
TC =
°C
TH =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
C=
J/°C
Second Determination of C
TC =
°C
TH =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
C=
J/°C
Average C =
J/°C
8B Experiment: Determination of Reaction Enthalpy
First Determination of ∆H
TI =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
heat =
J
∆H =
kJ/mol
Second Determination of ∆H
TI =
°C
Temperature inside calorimeter at 15 s
intervals:
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
°C
TF =
°C
heat =
J
∆H =
kJ/mol
Average ∆H =
kJ/mol
% error = [(|∆H| - |55.8|)/55.8] x 100 =
%
8C Exercises
1. TF = _______________°C.
3. Unknown metal = _______________ .
50
2. TF = _______________°C.
Lab Session 10, Experiment 9: Charles’ Law
The purpose of this experiment is to study the changes in the volume of a gas with changes in
temperature at constant pressure.
9A Experiment
1. Use a thoroughly dried 125 mL Erlenmeyer
flask for this experiment. If it is not dry,
rinse the flask with a small amount of
acetone or ethanol and place it upside-down
on a paper towel to dry.
2. Fit the flask with a one-hole rubber stopper
inserted with a short piece of dry glass
tubing. Add rubber tubing to the end of the
glass tube and assemble the apparatus as
shown in Figure 10.1, using the125 mL flask
and a 400 mL beaker. Be sure that the
stopper fits tightly in the flask, the glass
tubing fits tightly in the rubber stopper, and
the rubber tubing fits tightly on the glass
tubing. (The latter may require a hose clamp
or wire winding.) Leave a 1 cm gap between
the bottom of the flask and the beaker.
3. Pour water into the beaker until as much of
the flask is covered as possible.
4. Using a Bunsen burner, heat the water in the
Figure 10.1
beaker until it boils and then continue to heat
for 5 minutes. At this point determine the
temperature of the water by means of a 110 º C thermometer. Read the thermometer while
its bulb is immersed in the water and record the reading as entry (c) in the data table below.
Make sure the thermometer does not touch the beaker.
5. After the water in the beaker has been boiling for about 5 minutes, place a screw clamp (or a
pinch clamp) on the rubber tubing. While the water is still boiling, close the clamp in order
to make the flask airtight. Turn off the Bunsen burner by closing the gas valve at its base,
then turn off the gas at the bench gas cock.
6. Loosen the ring stand clamp, remove the flask from the beaker, and immerse it in the water
in the stoppered end sink nearest your work area. It should be cooled upside-down in the
sink for about 5 minutes, with the stopper and tubing completely immersed in water
throughout this period. At the end of this time, loosen the screw clamp while the end of the
tube is completely submerged and let water be forced into the flask.
7. Holding the rubber tubing closed with your fingers, remove the flask from the sink, and
determine the volume of the water in the flask by pouring it into a 100 mL graduated
cylinder. Record this volume as entry (a) in the data table.
51
8. Determine the temperature of the water in the sink and record it as entry (d) in the data table.
Your instructor may opt to prepare ice/water baths in the end sinks, in which case the
temperature should be close to 0°C. You should measure it with your thermometer and
record it anyway.
9. To determine the volume of gas (air) used in the beginning, it is necessary to accurately
determine the volume of the Erlenmeyer flask. Completely fill the flask with water and place
the rubber stopper (and tubing) in its previous position. Remove the stopper and measure the
volume of the water in the flask with a 100 mL graduated cylinder. Record this volume as
entry (b) in the data table.
(a) Volume of water forced into the flask
mL
(b) Initial volume of the air
(measured volume of the flask/stopper)
(c) Initial temperature of the air
(temperature of the hot water)
(d) Final temperature of the air
(temperature of the water in the sink)
(e) Final volume of the air [(b)–(a)]
(measured volume of the flask minus the volume of water
forced into the flask)
mL
°C
K
°C
K
mL
To summarize, a given volume of air was taken at the temperature of boiling water. The air was
cooled, causing it to contract. The new volume was determined experimentally.
9B Exercise
The accuracy of the experimental determination of the final volume (entry (e) in the data table)
can be checked by calculating this volume with Charles' Law. Change the temperature reading
into the absolute (Kelvin) scale and enter your data into the following formula:
Initial volume (mL) ×
Final Temperature (K)
Initial Temperature (K)
=
Final Volume (mL)
Your calculations:
×
=
Complete the following table.
(a) What was the experimental final volume?
mL
(b) According to Charles' Law, what should have been the final volume?
(c) What is the difference between the experimental and calculated
volumes?
(d) Assuming the calculated volume to be correct, determine the percent
error of the experiment.
52
mL
mL
%
Report Form 9: Charles' Law
Name__________________________________
Partner____________________Section #_____
9A Experiment
(a) Volume of water forced into the flask
mL
(b) Initial volume of the air
(measured volume of the flask/stopper)
(c) Initial temperature of the air
(temperature of the hot water)
(d) Final temperature of the air
(temperature of the water in the sink)
(e) Final volume of the air [(b)–(a)]
(measured volume of the flask minus the volume of water
forced into the flask)
mL
°C
K
°C
K
mL
9B Exercise
(a) What was the experimental final volume?
mL
(b) According to Charles' Law, what should have been the final volume?
(c) What is the difference between the experimental and calculated
volumes?
(d) Assuming the calculated volume to be correct, determine the percent
error of the experiment.
53
mL
mL
%
Lab Session 11, Experiment 10: Determination of the Molar
Mass of Oxygen
The molar mass (sometimes called the gram molar mass, or the molecular weight) of a gaseous
compound can be calculated if the following are known: the mass, volume, temperature, and
pressure of the gas. In this experiment, the molar mass of oxygen will be determined. Potassium
chlorate will be decomposed to produce oxygen. The mass of oxygen generated will equal the
difference between the mass of KClO3 before decomposition and the mass of the residue after
decomposition. The volume of the oxygen will be the volume of water displaced from a bottle
initially filled with water. The total gaseous pressure (Poxygen + Pwater) will be assumed to be
barometric pressure. The temperature of the oxygen will be assumed to be the same as that of
the water displaced.
10A Experiment
Record all measurements in the
data table provided.
1. Fill KClO3 into your large test
tube to a depth of 2-3 cm.
Add a small quantity of MnO2
and weigh the test tube with
contents.
2. Mix well. Mount the test tube
as shown in Figure 11.1.
Note: All connections must
be tight. This includes glass
tubes through stoppers,
stoppers in vessels, and rubber
tubing connections to glass
tubes. The latter may require
a hose clamp or wire winding.
Figure 11.1
3. Before heat is applied to the test tube, the assembly must be checked to make sure that water
will not be transferred from the bottle to the beaker by siphoning. Proceed with this check as
follows. Remove the test tube from the assembly by removing the rubber stopper from it.
Insert the glass tube extending from the stopper into the opening of a rubber bulb that you get
from the stockroom. Squeeze the rubber bulb gently to fill the delivery tube between the
bottle and the beaker with water. While the delivery tube is filled, close a pinch-clamp over
the connecting rubber tubing between the test tube and the bottle. Empty the beaker; a small
quantity of water will be transferred to the beaker when filling the delivery tube. Put the
large test tube back in place. Remove the pinch-clamp and observe for siphoning. If water
flows into the beaker, there is a leak around the stopper that must be corrected before
proceeding. If there is no evidence of siphoning, proceed to the next step.
54
4. Apply heat from your burner, slowly at first, to begin the decomposition of the chlorate.
After a little more than half the water in the bottle has been transferred to the beaker,
discontinue the heating.
5. Allow the test tube to cool to room temperature, then place a pinch-clamp over the rubber
tube. Remove and weigh the test tube. Measure the volume and temperature of the water in
the beaker, which is equivalent to the volume and temperature of oxygen generated.
6. Obtain the vapor pressure of water from the table below and calculate the pressure of dry
oxygen. You will have to convert the vapor pressure of water from Torr to atm.
[1 atm = 760 Torr]
Vapor Pressure of Water
ºC
18.0
18.2
18.4
18.6
18.8
19.0
19.2
19.4
19.6
19.8
Torr
15.477
15.673
15.871
16.071
16.272
16.477
16.685
16.894
17.105
17.319
ºC
20.0
20.2
20.4
20.6
20.8
21.0
21.2
21.4
21.6
21.8
Torr
17.535
17.753
17.974
18.197
18.422
18.650
18.880
19.113
19.349
19.587
ºC
22.0
22.2
22.4
22.6
22.8
23.0
23.2
23.4
23.6
23.8
Torr
19.827
20.070
20.316
20.565
20.815
21.068
21.324
21.583
21.845
22.110
ºC
24.0
24.2
24.4
24.6
24.8
25.0
25.2
25.4
25.6
25.8
Torr
22.377
22.648
22.922
23.198
23.476
23.756
24.039
24.326
24.617
24.912
ºC
26.0
26.2
26.4
26.6
26.8
27.0
27.2
27.4
27.6
27.8
Torr
25.209
25.509
25.812
26.117
26.426
26.739
27.055
27.374
27.696
28.021
ºC
28.0
28.2
28.4
28.6
28.8
29.0
29.2
29.4
29.6
29.8
Torr
28.349
28.680
29.015
29.354
29.697
30.043
30.392
30.745
31.102
31.461
7. Use the ideal gas law to calculate the gram-molar mass of oxygen. See the review of the ideal
gas law in the table below.
Review of the Ideal Gas Law
PV = nRT
M = (mass × RT) ÷ PV
PV = (mass ÷ M) × RT
where M is g/mol of gas
PVM = mass × RT
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
R = 0.08206 L atm mol
Mass of test tube and contents before heating
Mass of test tube and contents after heating
Mass of oxygen in flask [(a)–(b)]
Volume of water transferred = Volume of O2
Temperature of water = Temperature of O2
Barometric pressure = Patm
Vapor pressure of water
Pressure of dry oxygen [(f)–(g)]
Molar mass of oxygen
% error = [(|(i) – 32.00| / 32.00) x 100]
55
g
g
g
mL
ºC
Torr
Torr
Torr
g/mol
%
–1
–1
K
L
K
atm
atm
atm
10B Exercise
1. Why is it necessary to allow the test tube to cool before it is removed for weighing?
2. Why must the water delivery tube extend nearly to the bottom of the flask?
3. When water flow ceases, what can be assumed as the relationship between barometric
pressure (Patm) and the pressure inside the flask (Poxygen + Pwater)?
4. Is it necessary to decompose all the chlorate?
56
Name__________________________________
Partner____________________Section #_____
Report Form 10: Determination
of the Molar Mass of Oxygen
10A Experiment
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
Mass of test tube and contents before heating
Mass of test tube and contents after heating
Mass of oxygen in flask [(a)–(b)]
Volume of water transferred = Volume of O2
Temperature of water = Temperature of O2
Barometric pressure = Patm
Vapor pressure of water
Pressure of dry oxygen [(f)–(g)]
Molar mass of oxygen
% error = [(|(i) – 32.00| / 32.00) x 100]
g
g
g
mL
ºC
Torr
Torr
Torr
g/mol
%
10B Exercise
1. Why is it necessary to allow the test tube to cool before it is removed for weighing?
2. Why must the glass water delivery tube extend nearly to the bottom of the bottle?
3. When water flow ceases, what can be assumed as the relationship between barometric
pressure (Patm) and the pressure inside the flask (Poxygen + Pwater)?
4. Is it necessary to decompose all the chlorate?
57
L
K
atm
atm
atm
Lab Session 12, Experiment 11: Valence, the Combining
Capacity of Elements and Molecules
One molar mass of an element or compound that produces or reacts with 1.008 grams of
hydrogen has a valence of I. It follows that the same quantity producing or reacting with 2.016
grams of hydrogen is valence II, and producing or reacting with 3.024 grams of hydrogen is
valence III, etc.
In this experiment, the magnesium metal will be reacted with concentrated hydrochloric acid to
produce hydrogen gas according to the following chemical equation:
Mg (sol) + x HCl (aq) → MgCl x
(aq)
+
x
H2 (gas)
2
The hydrogen gas produced will be collected, its volume measured, and the mass in grams
calculated using the ideal gas law:
PV = nRT
When using this equation, the units of the gas constant (R) must be consistent with the units used
for the other quantities in the equation. The Handbook of Chemistry and Physics lists the value
of R as
R = 0.08206 L × atm × mol
–1
–1
×K
Thus, pressure (P) must be expressed in atmospheres, volume (V) in liters, and temperature (T)
in Kelvin.
11A Experiment
A sample of magnesium metal with an approximate mass of 0.24 grams will be reacted with
excess hydrochloric acid. The hydrogen gas produced will be collected, its volume measured,
and finally its mass in grams will be calculated.
Chemicals Required: 6M hydrochloric acid
Note: 12M HCl(aq) must be diluted.
Apparatus Required:
Magnesium metal,
ribbon or granules
125 mL Erlenmeyer flask
Metal or plastic trough
100 mL graduated cylinder
Thermometer
500 mL Florence flask
Glass tubing
Watch glass
Rubber tubing
58
The apparatus should be set up as
shown in Figure 12.1. Make sure
the bottom of the thistle tube
almost touches the bottom of the
Erlenmeyer flask. You may be
instructed to use the end sink
rather than a metal or plastic
trough. Be sure that all fittings
are tight. This includes the
fittings of the rubber stopper in
the Erlenmeyer flask, the glass
tube through the stopper, and the
rubber tubing to the glass tube.
(The latter may require a hose
clamp or wire winding.)
Figure 12.1
1. Record the exact mass of the magnesium as entry (a) in the data table.
2. Remove the two-hole stopper from the Erlenmeyer flask and add enough water the cover the
bottom of the thistle tube. If the magnesium is ribbon, form it into a ball or tight coil, and
place it in the flask. Reinsert the rubber stopper tightly.
3. Add 10 mL of 6M hydrochloric acid solution by pouring it into the thistle tube. Put a rubber
stopper in the top of the thistle tube to avoid back-splashing acid or escape of hydrogen gas.
4. After the reaction is complete, allow a few minutes for the flask to cool. Be sure to keep the
end of the rubber tubing above the water level in the Florence flask. The gas collected (H2
and water vapor) must be at the same pressure as the barometric pressure. To equalize
pressures, remove the gas delivery tube and raise or lower the Florence flask so that the water
levels in the trough (or end sink) and in the Florence flask are equal. (You may have to tilt
the flask to equalize levels. Be careful that the mouth of the Florence flask stays below the
water level in the trough.) Cover the Florence flask opening with a watch glass, invert the
Florence flask and place it on the desktop.
5. Using a graduated cylinder, measure the volume of water that must be added to the Florence
flask to fill it. It will be assumed that this is the volume of the gas collected. It should be
recorded as item (e) in the data table.
6. Measure the temperature of the water in the trough with your thermometer and record it as
item (b) below. It will be assumed that this is the temperature of the collected gas.
7. Record the barometric pressure as item (c) below. This is the total pressure of the hydrogen
and water vapor.
(a)
(b)
(c)
(d)
(e)
Mass of magnesium
Temperature of water in trough
Barometric pressure
Vapor pressure of water at the temperature recorded in (b)
Volume of gas collected = Volume of H2O added
59
g
ºC
Torr
Torr
mL
11B Exercise: Converting Data to the Correct Units
1. Calculate the Pressure
P = ___________atm
Pressure [atm] =
Barometric Pressure
–
Vapor Pressure of H2O
760 Torr/atm
2. Calculate the Volume
V = __________ L
Volume in L
=
Volume in mL
3. Calculate the Temperature
÷
1000
T = __________K
Absolute temperature in Kelvins
4. Calculate the mass of hydrogen
=
+
ºC
273.15
Mass = __________g
Calculating the grams of hydrogen produced using the ideal gas law:
PV = Mass of hydrogen produced ÷ molar mass of hydrogen × 0.08206 (R) × T
Rearranged:
Mass in grams of hydrogen produced = P × V × Molar Mass of H2
0.08206 (R) × T
5. From the generalized balanced equation: Mg + x HCl → MgClx + ½ x H2, we can calculate
how many grams of hydrogen should be produced from the number of grams of magnesium
reacted, as follows:
⎛ 1 mole Mg ⎞
⎟⎟
g hydrogen = (g Mg) ⎜⎜
⎝ 24.305 g Mg ⎠
⎛ (x/2) moles H 2 ⎞ ⎛ 2.016 g H 2 ⎞
x
⎟⎟ =
⎜⎜
⎟⎟ ⎜⎜
.
24.11
⎝ 1 mole Mg ⎠ ⎝ 1 mole H 2 ⎠
Rearrangement yields the following ratio.
g hydrogen
g Mg
=
x
. Thus, x =
24.11
⎛ g hydrogen ⎞
⎜
⎟ (24.11). x =
⎜ g
⎟
Mg
⎝
⎠
|x – 2|
X 100 = ________%
% error =
2
60
⎛ _________ g H 2
⎜
⎜ __________ g
Mg
⎝
⎞
⎟ (24.11) = ______ .
⎟
⎠
Give a balanced equation for the reaction of magnesium with hydrochloric acid:
6. Verification that HCl is the excess reagent by completing the following calculations:
g Mg
10 mL
1 mole Mg
24.3 g Mg
10 -3 L
1 mL
2 mole HCl
1 mole Mg
6.0 moles HCl
L
=
____________ moles HCL needed
=
____________ moles HCL added
________ moles HCl added – ________ moles HCl needed = ________ moles of HCl unreacted
11C Alternate Calculations
1.
Moles of H2
= PV = (____________)(____________) = __________mol.
RT
(____________)(____________)
2.
Moles of Mg
3.
Mole ratio of H2 to Mg
4.
xexp
5.
% error
= gMg
MMg
= 2 × ratio
=
=
xexp – 2.00
2.00
=
(____________)
(____________)
=
Moles of H2
Moles of Mg
=
2 × (____________)
=
= __________mol.
(____________) =
(____________)
=
__________.
(____________) – 2.00
2.00
=
__________%.
11D Exercise: Nomenclature
Give the formula of each of the following.
calcium sulfide
Iron (II) chloride
sodium sulfate
nickel (II) nitrate
calcium phosphate
magnesium hydroxide
lithium oxide
Tin (IV) bromide
61
__________.
Name__________________________________
Partner____________________Section #_____
Report Form 11: The Combining
Capacity of Elements and
Molecules
11A Experiment
(a)
(b)
(c)
(d)
(e)
Mass of magnesium
Temperature of water in trough
Barometric pressure
Vapor pressure of water at the temperature recorded in (b)
Volume of gas collected = Volume of H2O added
g
ºC
Torr
Torr
mL
11B Exercise: Converting Data to the Correct Units
1. Calculate the Pressure
P = ________atm
2. Calculate the Volume
V = __________L
3. Calculate the Temperature
T = __________K
4. Calculate the Mass of hydrogen
⎛ __________ g H 2
5. x = ⎜
⎜ __________ g
Mg
⎝
Mass = __________g
⎞
⎟ (24.11) = ________ .
⎟
⎠
% error = ________.
Give a balanced equation for the reaction of magnesium with hydrochloric acid:
____________________________________________________________________
6. Verification that HCl is the excess reagent by completing the following calculations:
________ moles HCl added – ________ moles HCl needed = ________ moles of HCl unreacted
11C Alternate Calculations
1. Moles of H2 = ______________mol.
4. xexp = 2 × ratio = ______________.
2. Moles of Mg = ______________mol.
5. % error = ______________%.
3. Mole ratio of H2 to Mg = ______________.
11D Exercise: Nomenclature
calcium sulfide
Iron (II) chloride
sodium sulfate
nickel (II) nitrate
calcium phosphate
magnesium hydroxide
lithium oxide
Tin (IV) bromide
62
Lab Session 13, Experiment 12: Copper Analysis by
Complexometric Titration
A quantitative analysis of copper in a soluble copper salt will by performed by complexometric
titration. The complexing agent will be ethylenediaminetetraacetic acid (EDTA) in the form of
–1
its disodium dihydrate salt (Na2C10H18N2O10), with a molar mass of 372.24 g mol . Since
EDTA forms complexes with many metal ions, this particular method can only be used in the
2+
2+
absence of such ions as Ca , Ni , etc. The reaction of complexation is:
2+
Cu
(aq)
+ (EDTA)
2–
(aq)
→ Cu(EDTA)
2–
+
(aq)
+ 2H
(aq)
2+
The stoichiometry is one metal cation to one EDTA anion. However, for Cu (since it has lost 2
–1
2–
electrons), the equivalent mass is 63.546/2 = 31.773 g eq , and since (EDTA) is a dianion its
–1
equivalent mass is 372.24/2 = 186.12 g eq (for the disodium dihydrate salt). The equation
above represents two equivalents reacting with two equivalents. The complex dianion is formed
+
with the release of two moles of H from EDTA, with the indicator being released from the
copper ion.
The complex dianion has the structure shown
in Figure 13.1. Note that the anion
completely surrounds the cation, forming six
coordinate covalent bonds to copper and a
very stable complex. The bonding to the
copper ion is nearly octahedral.
The indicator used for the titration is called
murexide. This indicator is highly colored
and will complex with the copper ion to give
a different colored species. During the
2–
titration, the EDTA forms a more stable
Figure 13.1
complex and frees the indicator, which then
Structure
of
Cu2+— EDTA dianion
displays its original color. The appearance of
the free indicator means that all metal ions
2–
have been complexed by EDTA , which
signals the end point. At the end point, the following equation applies:
NEDTAVEDTA = NCu(II)VCu(II) = #eq Cu(II), if V is given in L
= #meq Cu(II), if V is given in mL
The mass of Cu equals (#eq Cu(II)) × (equivalent mass of Cu(II)), and
mass Cu(II)
mass Cu(II) salt
63
× 100 = % Cu
12A Experiment
1. Rinse your burette and fill it with standardized Na2EDTA•2H2O solution
(7.445 g Na2EDTA • 2 H2O per liter of water).
2. Weigh accurately three approximately 0.1 g samples of the copper salt.
3. Dissolve each sample in 50 mL of de-ionized water.
4. Add exactly the same amount of indicator to each sample, three drops to start off with. If the
indicator solution is not strong enough, add more but always the same for all samples. (The
indicator's concentration should be 100 mg/100 mL H2O)
5. Titrate each sample with the standardized EDTA. The light yellow solution turns green near
the end point, then suddenly purplish blue at the end point. This end point is fairly hard to
see, so put a white sheet of paper under your beaker and watch carefully. The distinctly
purplish hue, due to free murexide, is the key to observing the end point.
6. For each titration, calculate the number of equivalents or (milliequivalents) of Cu(II) found.
7. For each titrated sample, calculate the mass of copper in that sample.
8. For each titration, calculate the % copper content in the sample, then average them.
Sample 1
(a)
(b)
(c)
(d)
(e)
(f)
Sample 2
Normality of EDTA
grams of Cu(II) sample
mL of EDTA solution
eq (or meq) of Cu(II)
mass of copper
% copper content
Sample 3
Average
% copper
content
12B Exercise
In the experiment on hydrates, we found that copper sulfate was a hydrate which contained
2–
36.1% by mass water. Since the only other component is the sulfate ion, SO4 , we can now
determine the complete formula of copper sulfate.
NOTE Solutions preparation:
Either weigh the EDTA analytically or standardize the solution. Label the bottles with the
normality of EDTA. Use deionized water. About 1 liter will be used by 20 students, 10 groups.
Make these solutions up fresh, including the murexide solution.
64
Report Form 12: Copper Analysis Name ________________________________
Partner___________________Section #_____
by Complexometric Titration
12A Experiment
Sample 1
(g)
(h)
(i)
(j)
(k)
(l)
Sample 2
Normality of EDTA
grams of Cu(II) sample
mL of EDTA solution
eq (or meq) of Cu(II)
mass of copper
% copper content
Sample 3
Average
% copper
content
12B Exercise
Complete formula of copper sulfate.
65